Why Do Periodic Trends Vary in Groups and Periods?
Understanding how properties of elements vary down a group and across a period is essential for mastering periodic trends in Chemistry, a high-yield topic for JEE and other exams. Periodic properties like atomic size, ionisation enthalpy, electron affinity, and oxide character display clear patterns on the periodic table, helping in problem-solving and prediction of chemical behavior.
1. What is Periodicity in Periodic Properties?
Periodicity is the repetition of certain physical and chemical properties of elements at regular intervals when arranged by increasing atomic number. This happens due to the recurring similar electronic configurations in the outer shells as you move across periods or down groups. For example, all alkali metals (Group 1) have a single electron in their outer shell, so they share common characteristics, such as high reactivity and tendency to form +1 cations.
- Atomic Size and how it changes on the periodic table
- Difference Between Elements and Atoms (foundation overview)
- Group 16 Elements, Group 17 Elements, Group 18 Elements examples of periodicity in groups
2. Variations in Key Periodic Properties
Let’s break down how vital periodic properties evolve both down a group (top to bottom) and across a period (left to right):
| Property | Trend Down a Group | Trend Across a Period |
|---|---|---|
| Atomic Size (Radius) | Increases (new shells added) |
Decreases (stronger nuclear pull) |
| Ionisation Enthalpy (Energy to remove electron) |
Decreases (valence electron farther from nucleus, more shielding) |
Increases (nucleus holds electrons more tightly) |
| Electron Affinity (Energy released on gaining electron) |
Becomes less negative (added electron farther from nucleus) |
Becomes more negative (smaller size, higher nuclear charge, easier to attract electrons) |
| Electronegativity | Decreases (less tendency to attract electrons) |
Increases (non-metals on right attract electrons more) |
| Nature of Oxides | Basic character increases (left side) Acidic character increases (right side) |
Basic oxides → Amphoteric → Acidic oxides (across period) |
3. Explanation and Examples
Atomic Size (Radius): Going down a group, each new element has an extra shell of electrons, so atoms get larger. Across a period, electrons are added to the same shell while the nuclear charge increases, pulling electrons closer and reducing size.
Ionisation Enthalpy: It’s easier to remove an outer electron from a large atom (down a group), so ionisation enthalpy falls. Across a period, the increasing nuclear charge makes it harder to remove electrons, so ionisation enthalpy rises.
Example: First ionisation energies: Li < Be < B < C < N < O < F < Ne (across Period 2).
Electron Affinity: Higher (more negative) when a new electron is strongly attracted to the nucleus. Halogens have the highest electron affinity. As we move down, the added electron is farther from the nucleus, so the attraction weakens.
Exception: Oxygen and fluorine have less negative electron affinities than S and Cl due to small size and increased electron-electron repulsion in their compact second shells.
Nature of Oxides: Elements on the left (s-block) form basic oxides, those in the middle (transition/metalloids) have amphoteric or neutral oxides, and those on the right (p-block non-metals) form acidic oxides.
Example: Third period oxides transition:
Na2O (strongly basic) → MgO (basic) → Al2O3 (amphoteric) → SiO2 (weak acidic) → P2O5 (acidic) → SO3 (strongly acidic) → Cl2O7 (extremely acidic).
4. Exam Notes and Tips
- Atomic size increases down a group but decreases across a period.
- Ionisation enthalpy shows opposite behavior (decreases down, increases across).
- Electron affinity becomes more negative across a period, trends down a group vary with exceptions.
- Remember oxide trend: basic → amphoteric → acidic across a period.
- Real world: Periodic property trends explain metal reactivity, salt formation, and environmental chemistry.
- JEE questions often ask for order of atomic/ionic size, ionisation energies, or nature of oxides.
5. Frequently Asked Example Questions
- Arrange Na, K, and Cs in order of increasing atomic size.
Answer: Na < K < Cs - Which element in Group 1 has the highest ionisation enthalpy?
Answer: Lithium (Li) - Why does Cl have higher electron affinity than F?
Answer: Less electron-electron repulsion in Cl's larger 3p orbital compared to F's compact 2p orbital. - Classify the nature of SO3 oxide:
Answer: Strongly acidic oxide
6. Summary Table of Key Trends
| Trend | Down a Group | Across a Period |
|---|---|---|
| Atomic Radius | Increases | Decreases |
| Ionisation Energy | Decreases | Increases |
| Electron Affinity | Becomes less negative | Becomes more negative |
| Electronegativity | Decreases | Increases |
7. Conclusion
The modern periodic table’s genius lies in how it groups elements so their properties vary in predictable and examinable ways. These trends – in atomic size, ionisation energy, electron affinity, electronegativity, and nature of oxides – are at the heart of chemical reasoning, reactivity, and periodicity. Periodicity is crucial for predicting properties of unfamiliar elements and for tackling multiple-choice and conceptual problems in JEE and board exams.
- Next, solidify your grasp by reviewing Atomic Size and group trends in Group 16 Elements and others.
- For a foundation refresher, visit Difference Between Elements and Atoms.
FAQs on How Properties Change Down a Group and Across a Period
1. What is the trend in atomic radius down a group and across a period?
Atomic radius increases down a group and decreases across a period in the periodic table.
- Down a group: New electron shells are added, so atoms get larger.
- Across a period: Electrons are added to the same shell, but increased nuclear charge pulls them closer, reducing atomic size.
2. How does ionization enthalpy vary down a group and along a period?
Ionization enthalpy decreases down a group and increases along a period.
- Down a group: Electrons are farther from the nucleus, so less energy is needed to remove them.
- Across a period: Higher nuclear charge makes it harder to remove an electron, so ionization enthalpy increases.
3. Why does electronegativity increase across a period?
Electronegativity increases across a period because the nuclear charge increases, attracting the bonding electrons more strongly.
- Atoms have greater tendency to attract electrons in a bond
- Atomic size decreases, allowing the nucleus to pull electrons closer
4. Describe the trend of metallic character in the periodic table.
Metallic character increases down a group and decreases across a period.
- Down a group: Atoms lose electrons more easily, becoming more metallic.
- Across a period: Atoms hold electrons more tightly, so they are less metallic.
5. What is the trend of electron affinity down a group and along a period?
Electron affinity decreases down a group and generally increases across a period.
- Down a group: Larger atoms have less attraction for incoming electrons.
- Across a period: Higher nuclear charge attracts added electrons more strongly, raising electron affinity.
6. Which property increases down a group in the periodic table?
Properties like atomic radius and metallic character increase as you move down a group.
- Additional electron shells make the atom bigger.
- Atoms lose electrons more easily, increasing metallic character.
7. Why does atomic size decrease across a period from left to right?
Atomic size decreases across a period due to increase in nuclear charge, which pulls electrons closer to the nucleus.
- No new electron shells are added across a period.
- Greater pull leads to a smaller atomic radius.
8. Explain the variation of non-metallic character along a period and down a group.
Non-metallic character increases across a period and decreases down a group.
- Across a period: Elements gain electrons more easily, becoming non-metals.
- Down a group: Elements tend to lose electrons, showing more metallic character.
9. What happens to reactivity of alkali metals down the group?
The reactivity of alkali metals increases down the group.
- Atomic size increases, so the outer electron is removed more easily.
- Lower ionization energy means higher reactivity.
10. How do chemical properties of elements show periodicity?
Chemical properties of elements change in a regular pattern, called periodicity, due to repeating outer electronic configurations.
- Atomic radius, ionization enthalpy, and electronegativity show recurring trends.
- These trends repeat across periods and down groups in the periodic table.
